Electrochemical equilibrium diagrams

Pourbaix plotted electrochemical equilibrium diagrams of metals in water as a function of the potential E with respect to the hydrogen electrode, and as a function of pH (Figure B.1.10). Several domains can be identified in these diagrams corrosion, passivation and immunity (see Section B.1.6). 

Frankenthal R P and Kruger J (eds) 1984 Equilibrium Diagrams of Localized Corrosion Proc. vol 84-9 (Pennington, NJ Electrochemical Society)... 

Potential-pH Equilibrium Diagram (Pourbaix Diagram) diagram of the equilibrium potentials of electrochemical reactions as a function of the pH of the solution. The diagram shows the phases that are thermodynamically stable when a metal reacts with water or an aqueous solution of specified ions. 

The analysis of thermodynamic data obeying chemical and electrochemical equilibrium is essential in understanding the reactivity of a system to be used for deposition/synthesis of a desired phase prior to moving to experiment and/or implementing complementary kinetic analysis tools. Theoretical and (quasi-)equilibrium data can be summarized in Pourbaix (potential-pH) diagrams, which may provide a comprehensive picture of the electrochemical solution growth system in terms of variables and reaction possibilities under different conditions of pH, redox potential, and/or concentrations of dissolved and electroactive substances. 

Figure 2.14 Electrochemical phase diagram for chalcopyrite with elemental sulphur as metastable phase. Equilibrium lines (solid lines) correspond to dissolved species at 10 mol/L. Plotted points show the upper and lower limit potential of collectorless flotation of chalcopyrite reported from Sun (1990), Feng (1989) and Trahar (1984)...

Electrochemical phase diagrams have been used to investigate the collector water mineral system in which the experimental potential for flotation is compared with thermodynamic equilibriums for reactions in mineral/oxygen/collector system to... 

From the Nernst equation, proton-coupled electron addition leads to a - 59 mV shift in potential per pH unit. Figure 15 shows the behavior of titanium dioxide and Figure 16 shows the behavior of tin oxide. The plots comprise Pour-baix diagrams for these materials. The breaks observed at extreme pHs with 2 define pAfa s for TitvO(OH) and TimO(OH), with the relevant electrochemical equilibrium at less extreme pH or 0 values [ — 8 < pH ( 0) < 23] described by [78]... 

The skills developed to produce the equilibrium diagram Figure A.l, are now applied anew. Neither hydrogen nor carbon monoxide occur as free substances in nature, where they are immediately oxidized. They must be made and stored, at thermodynamic and economic cost. The reversible thermodynamics are assessed below, using as the basis of calculation a notional, electrochemical, equilibrium, steam reformer. Figure A.4, for comparison with the alternative practical and irreversible combustion-driven reformers. 

Figure 24. Free energy diagrams corresponding to the heuristic model for electrochemical reactions presented in section 7.4. The upper diagram represents the situation at electrochemical equilibrium the lower diagram represents the situation where U > f/ (nett anodic current flow).

The description of corrosion kinetics in electrochemical terms is based on the use of potential-current diagrams and a consideration of polarization effects. The equilibrium or reversible potentials Involved in the construction of equilibrium diagrams assume that there is no net transfer of charge (the anodic and cathodic currents are approximately zero). When the current flow is not zero, the anodic and cathodic potentials of the corrosion cell differ from their equilibrium values the anodic potential becomes, more positive, and the cathodic potential becomes more negative. The voltage difference, or polarization, can be due to cell resistance (resistance polarization) to the depletion of a reactant or the build-up of a product at an electrode surface (concentration polarization) or to a slow step in an electrode reaction (activation polarization). 

L. D. Burke and R. A. Scanned, Equilibrium Diagrams Localized Corrosion (Proceeding of an International Symposium Honoring Marcel Pourbaix on his Eightieth Birthday), Ed. by R. P. Frankenthal and J. Kruger, The Electrochemical Society, Pennington, New Jersey, 1984, pp. 135-147. 

Frankenthal, R. P. and Kruger, J., Equilibrium Diagrams—Localized Corrosion, Electrochemical Society, Pennington, NJ, 1984. 

Dissolution Potential of Aluminium Electrochemical Equilibrium (Pourbaix) Diagrams The Electrochemical Behaviour of Aluminium Aluminium as a Passive Metal... 

The hydrogen ion concentration - expressed by pH - is an important parameter for calculating electrochemical equilibrium potentials. The Pourbaix diagram maps in a clear way the influence of the pH value on complicated electrochemical equilibria. The diagram form, which is named after the Belgian corrosion researcher Marcel Pourbaix, was developed in the 1940s. 

The Pourbaix diagram specifies electrochemical equilibrium curves for metals and metal oxides in a voltage vs. pH coordinate system. These curves delimit areas where the metal is immune, where the metal is passivated, and areas where the metal is corrosion active at equilibrium conditions. These equilibrium curves can usually be calculated from the thermodynamic data of the substances areas with passivation or corrosion are determined by tests and from practical experience. 

Electrochemical potential differences can arise between two electrodes of the same metal if the electrodes are in contact with electrolytes of different composition. This is, for example one of the reasons why deposited salt particles on a moist metal surface can induce local galvanic corrosion cells on the surface. Approximate from the Nernst equation the electrochemical equilibrium potential AV = V2 — Vi (volt) at 25 °C for a Zn-Zn electrode pair with the following cell diagram... 

Ladder diagrams can also be used to evaluate equilibrium reactions in redox systems. Figure 6.9 shows a typical ladder diagram for two half-reactions in which the scale is the electrochemical potential, E. Areas of predominance are defined by the Nernst equation. Using the Fe +/Fe + half-reaction as an example, we write... 

Nickel occupies an intermediate position in the electrochemical series Ni2 /Ni = -0-227 V, so that it is more noble than Zn and Fe but less noble than Sn, Pb and Cu. Figure 4.21 shows a revised potential-pH equilibrium (Pourbaix) diagram for the Ni-H O system at 25°C. The existence of the higher anhydrous oxides Nij04, NijO, and NiOj shown in an earlier diagram appears doubtful in aqueous systems in the absence of positive identification of such species. It is seen that ... 

The general thermodynamic treatment of binary systems which involve the incorporation of an electroactive species into a solid alloy electrode under the assumption of complete equilibrium was presented by Weppner and Huggins [19-21], Under these conditions the Gibbs Phase Rule specifies that the electrochemical potential varies with composition in the single-phase regions of a binary phase diagram, and is composition-independent in two-phase regions if the temperature and total pressure are kept constant. 

In practice, for a ternary system, the decomposition voltage of the solid electrolyte may be readily measured with the help of a galvanic cell which makes use of the solid electrolyte under investigation and the adjacent equilibrium phase in the phase diagram as an electrode. A convenient technique is the formation of these phases electrochemically by decomposition of the electrolyte. The sample is polarized between a reversible electrode and an inert electrode such as Pt or Mo in the case of a lithium ion conductor, in the same direction as in polarization experiments. The... 

Chapters 7 to 9 apply the thermodynamic relationships to mixtures, to phase equilibria, and to chemical equilibrium. In Chapter 7, both nonelectrolyte and electrolyte solutions are described, including the properties of ideal mixtures. The Debye-Hiickel theory is developed and applied to the electrolyte solutions. Thermal properties and osmotic pressure are also described. In Chapter 8, the principles of phase equilibria of pure substances and of mixtures are presented. The phase rule, Clapeyron equation, and phase diagrams are used extensively in the description of representative systems. Chapter 9 uses thermodynamics to describe chemical equilibrium. The equilibrium constant and its relationship to pressure, temperature, and activity is developed, as are the basic equations that apply to electrochemical cells. Examples are given that demonstrate the use of thermodynamics in predicting equilibrium conditions and cell voltages. 

Figure 29.4 shows an example, the energy diagram of a cell where n-type cadmium sulfide CdS is used as a photoanode, a metal that is corrosion resistant and catalytically active is used as the (dark) cathode, and an alkaline solution with S and S2 ions between which the redox equilibrium S + 2e 2S exists is used as the electrolyte. In this system, equilibrium is practically established, not only at the metal-solution interface but also at the semiconductor-solution interface. Hence, in the dark, the electrochemical potentials of the electrons in all three phases are identical. 

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